A stable internal environment requires a precisely regulated acidity level, measured on the pH scale by the concentration of hydrogen ions (\(\text{H}^+\)). In the human body, the pH of blood must remain within a narrow range of 7.35 to 7.45; significant deviation can be life-threatening.
This strict control is necessary because the shape and function of proteins, particularly enzymes that catalyze cellular reactions, are highly sensitive to hydrogen ion concentration. If the environment becomes too acidic or too alkaline, these proteins lose their functional structure, halting cellular processes. Buffer systems are the body’s immediate, first line of defense against the constant influx of acids produced by metabolism, preventing sudden shifts in pH.
The Chemical Mechanism of Buffering
A chemical buffer system is a solution that resists changes in pH when a small amount of acid or base is added. Buffers consist of a weak acid and its corresponding conjugate base, which exist in equilibrium. The weak acid does not fully dissociate in water, allowing both components to remain present in high concentrations to absorb changes.
When a strong acid is introduced, it releases hydrogen ions (\(\text{H}^+\)) into the solution. The buffer’s conjugate base immediately reacts with these added \(\text{H}^+\) ions, binding them to form the original weak acid. This action prevents the free hydrogen ions from accumulating and causing a sharp drop in pH.
Conversely, if a strong base is added, it releases hydroxide ions (\(\text{OH}^-\)), which would cause the pH to rise rapidly. The weak acid neutralizes the added base by donating a hydrogen ion to the \(\text{OH}^-\) to form water. This converts the strong base into a weak conjugate base, minimizing the increase in alkalinity.
The effectiveness of any buffer is determined by its buffer capacity—the amount of acid or base it can absorb before the pH changes significantly. The system works most efficiently when the concentrations of the weak acid and its conjugate base are roughly equal. Once one component is depleted, the buffer capacity is exceeded, and the pH will fluctuate dramatically with any further addition of acid or base.
Major Buffer Systems in the Human Body
The human body employs three major buffer systems, each located to manage pH in different fluid compartments. The Bicarbonate Buffer System is the most important for systemic pH control, operating predominantly in the extracellular fluid, including blood plasma. It consists of carbonic acid (\(\text{H}_2\text{CO}_3\)) and bicarbonate ions (\(\text{HCO}_3^-\)).
This system is unique because its components are tightly linked to the respiratory and renal systems, allowing for constant regulation. Carbonic acid is formed when carbon dioxide (\(\text{CO}_2\)) reacts with water, meaning changes in breathing rate can rapidly alter the amount of acid present. This dynamic control makes the bicarbonate system powerful for maintaining the blood’s narrow pH range.
The Phosphate Buffer System is less concentrated in the blood but is highly effective in two locations: the intracellular fluid and the renal tubules. Its components are dihydrogen phosphate (\(\text{H}_2\text{PO}_4^-\)) and hydrogen phosphate (\(\text{HPO}_4^{2-}\)), which readily absorb or release hydrogen ions. In the kidney, this system is important for allowing the excretion of excess hydrogen ions into the urine.
The Protein Buffer System is the most abundant buffering mechanism inside cells, accounting for a significant portion of the body’s total chemical buffering capacity. Proteins are composed of amino acids, which contain carboxyl (\(\text{-COOH}\)) and amino (\(\text{-NH}_2\)) groups. The carboxyl group can release \(\text{H}^+\) when alkaline, while the amino group can bind \(\text{H}^+\) when acidic. Hemoglobin, the protein inside red blood cells, is a key example, buffering the hydrogen ions generated when carbon dioxide is transported from the tissues to the lungs.
Buffers and Physiological Stability
While buffer systems are the immediate chemical line of defense, they do not eliminate acids or bases from the body; they merely bind them temporarily. If metabolic processes generate excessive acid, buffers can become overwhelmed, leading to acidosis (pH drops below 7.35). Conversely, a loss of acid or an excess of base results in alkalosis (high pH).
To support and restore chemical buffers, the body relies on the respiratory and renal systems for long-term regulation. The respiratory system offers a rapid, temporary solution by controlling the exhalation of \(\text{CO}_2\). If the blood becomes too acidic, increasing the breathing rate quickly lowers the \(\text{CO}_2\) level, reducing the acid component of the bicarbonate buffer system within minutes.
The renal system provides the slower, more permanent means of compensation, taking hours to days to fully respond. The kidneys adjust the pH balance by controlling the reabsorption of bicarbonate ions and the excretion of fixed acids, such as phosphate, into the urine. This dual-organ support ensures that buffer systems are constantly replenished and protected, maintaining the acid-base balance.